INTERACTION OF HYDROGEN PEROXIDE AND HYPOCHLOROUS ACID
June, 1947
1COSTRIBUTIOSFROM
THE
1509
CHEMICAL LABORATORIES O F THE UXIVERSITY O F CALIFORNIA]
The Interaction of Hydrogen Peroxide and Hypochlorous Acid in Acidic Solutions Containing Chloride Ion BY ROBERT E. CONNICK Introduction The reaction between hydrogen peroxide and chlorine in hydrochloric acid solutions proceeds according to the equation HpOn
+ Clz
0 3
f 2Hf
+ 2C1-
(1)
The rate of this reaction has been studied by Makower and Bray.' In hydrochloric acid solutions of greater than one molal concentration they found the rate law to be
- d(HnO?) _ _ - k, dt
(HsO,)(Cl?) =
E (HzO,)(HOCl)(H+)(Cl-) k8
(8)
where K is the hydrolysis constant for chlorine. The present study was undertaken to investigate the validity of this rate law expression a t low concentrations of H + and C1-, and hence, to test the proposed mechanism. Makower's2 preliminary experiments had indicated that the rate law under these conditions was complex.
Experimental They accounted for this equation by assuming that the chlorine was in rapid equilibrium with its hydrolysis products and that the rate determining step was the reaction between hydrogen peroxide and hypochlorous acid, i.a. HOCl
H?O Clp H + + C1- f HOCl + HpOz +HpO f f H + f C10 2
(3) (4)
At hydrochloric acid concentrations below one molal, Makower and Bray found that the rate was less than that to be expected on the basis of the above equation. Later, in experiments at quite low hydrochloric acid concentration, Makower2p3 showed that the specific rate of the reaction was much less than for strong acid solutions, indicating that the above mechanism cannot be correct. I t was suggested to me by Professor Bray in 1939 that the experimental results might be explained by the following mechanism H?O1 4- Cl?
k, H + f C1- f HOOCl
(5)
k I, HOOCl
+0 2 + H kc
+ C1-
(6)
The resulting rate law is
For high values of the product of the hydrogen ion and chloride ion concentrations, this rate law reduces to that found experimentally by Makower and Bray, equation (2). At lower concentrations of H+ and C1-, the rate should decrease below that indicated by equation ( 2 ) as k, becomes significant, and a t sufficiently low concentrations of H+ and C1- should approach the expression (1) B. Makower and W. C. Bray, THISJ O U R N A L , 56,4765 (1933). For results on the catalytic decomposition of hydrogen peroxide by the simultaneous occurrence of this reaction and the oxidation of C1- t o CIP,by hydrogen peroxide, see Livingston and Bray, ibid., 47, 2069 (1926). and Maass and Hiebert, ibid., 46, 290 (1921). (la) Parentheses indicate the concentration of the enclosed substance in moles per liter of solution. (2) B. hlakower, Thesis, University of California, 1932. (3) W. C. Bray, Cltem. Reo., 10, 175 (1932).
Reagents.-Considerable care was taken t o exclude impurities because of their possible catalytic effect on the reaction. Conductivity water, prepared by redistillation of distilled water from alkaline permanganate solution in a tin still, was used throughout. Hydrogen peroxide solutions were made by dilution of Merck Superoxol, free from preservative. Solutions prepared from different bottles of this reagent showed no experimental differences in reaction rates. Chlorine solutions were prepared by heating manganese dioxide with 4 &I hydrochloric acid and passing the evolved gas through two bubblers filled with water and then into conductivity water. h-0 differences in behavior were ever detected among the many solutions prepared in this way. Hypochlorous acid solutions, which were used in a few experiments, were prepared by adding excess mercuric oxide t o a chlorine solution and distilling off the hypochlorous acid. These solutions showed no detectable decomposition during a storage period of two months. Solutions of sodium perchlorate were prepared by neutralizing pure 60% perchloric acid with sodium bicarbonate. Solutions of sodium chloride. were prepared from the C. P . salt. Analyses.-The reaction was followed by measuring the sum of the hydrogen peroxide and hypochlorous acid (or chlorine) concentrations. This was determined by adding samples from the reaction mixture to solutions made from 1 ml. of 1 ill potassium iodide, 1 ml. of 1 M sulfuric acid, 1 drop of 0.03 M ammonium molybdate and 20 ml. of water, and titrating the liberated iodine with standardized sodium thiosulfate. The very rapid reduction of the hypochlorous acid (or chlorine) by iodide stopped the reaction. The ammonium molybdate was used as a catalyst* for the relatively slow reduction of hydrogen peroxide by iodide ion. Samples were allowed t o stand ten minutes in the dark before titration with the thiosulfate to insure completion of the reaction. In a few experiments the hypochlorous acid concentration alone was obtained by, adding the sample t o a potassium iodide solution a t 0 , containing low acid and no molybdate, and titrating quickly with thiosulfate. By this procedure the reaction of hydrogen peroxide with iodide ion is made negligible. The results were concordant with thosc based on the other method of analysis. Experimental Procedure.-Measured volumes of solutions of hypochlorous acid and hydrogen peroxide of known concentration were placed in separate glass-stoppered flasks and were brought t o the reaction temperature by immersion in a thermostat for over an hour. Added hydrogen ion and chloride ion were contained in either or both of the flasks. The reaction was started by pouring (4) I. M . Kolthoff and N. H. Furman, "Volumetric Analysis," Vol. 11, John Wiley and Sons, New York, N. Y., 1929, p. 393.
ROBERT E.CONNICK
1510
Expt.
75 VI1 90 VI 96 IV 9 v 6V 92 VI 102 IV 107 VI 110 VI
Hi02
x
10‘
17.83 11.66 9.61 9.33 9.34 11.52 9.60 2.688 2.681
HOCl f Clr X 10‘
30.25 5.27 7.15 6.39 6.37 5.02 6.80 1.485 2,294
TABLE I DATAFOR kr” AT 25’ Initial concentrations in moles per liter. HOCl CI H+ ( H i ) \$\-I
x
104
29.87 4.48 6.03 5.40 5.34 4.04 2.92 0.1133 0.1698
x
104
29. S i 105.0 69.4 1.89.6 69.7 527 70.2 919 036
the contents of one flask into the other, zero time being taken a t the start of the mixing. For rapid reactions this “zero” time was not used. Immediately after mixing, the reaction flask was withdrawn from the bath, shaken vigorously, and immediately replaced. The mixing and shaking required from ten t o fifteen seconds. No effects attributable t o variations in the order or manner of mixing were ever observed. Samples were removed a t convenient time intervals with a 20-1111. pipet, the time being recorded when the pipet had half emptied into the potassium iodide solution. The pipet used for the more rapidly changing reaction mixtures had a delivery time of less than five seconds. In all experiments sodium perchlorate was used t o adjust the ionic strength to the desired value. Calculation of Reaction Rates.-Titers of the different samples were plotted against time and a smooth curve drawn through the points. Slopes were measured a t various points on the curve and these values, when multiplied by a suitable factor, gave the rate of disappearance of hydrogen peroxide, -d(HZOz)/dt. Rate constants were calculated using these figures together with the concentrations of the reacting species a t the time a t which the slope was measured. All concentrations have been expressed in moles per liter of solution a t the temperature of the experiment and time in minutes. In runs at high concentrations of hydrochloric acid there was appreciable loss of chlorine because of its volatility. This was assumed to have occurred principally during the equilibration and mixing period. Consequently, for such experiments the initial concentrations of chlorine were calculated from the final titer of the reagent present in excess and the initial concentration of hydrogen peroxide. For cases in which the chlorine concentration was small relative t o the hypochlorous acid concentration, this correction was small, but amounted t o as much as 15% for cases in which chlorine was the principal species. Hence, the results of these latter experiments are less accurate than the others. A comparison of the final titer of the reaction mixture with the difference in the initial concentrations of hydrogen peroxide and hypochlorous acid (or chlorine) showed that the assumed stoichiometry was closely followed for experiments in which loss of chlorine was not to be expected. MakowerZhad previously shown that the volume of oxygen liberated under these conditions is in accord with this stoichiometry.
Results and Discussion The kq Law.-When the product of the hydrogen ion and chloride ion concentrations lies within the limits lo-’ to the experimentally determined rate law is
- d(HlOz)/dt
Vol. 69
= k~(HOCl)(HiOz)(H+)(Cl-)
(9)
Table I is a summary of the runs used in establishing this law. The headings are self-explanatory except €or the last two columns which are re-
x
10’
29.87 118.0 186 71.2 205 33.7 1393 919 936
0.885 12.39 12.9 13.50 14.29 17.76 97.8 845 876
x
k4
10-7
2.00 1.90 1.71 1.75 1.70 2.04 1.74 1.50 1.55
P
0.0958 .lo3 .289 .158 .160 .149 .146 .0920 .0937
x
kdo
10-
3.11 3.01 2.95 2.89 2.81 3.36 2.86 2.3 2.4
spectively the ionic strength, p, and the values of
k4 extrapolated to zero ionic strength (to be discussed later). Each value of k4 represents the average of several values for that experiment calculated a t different times. In calculating the true concentration of hypochlorous acid, i t was necessary to use the chlorine hydrolysis constant. The extensive measurements of Jakowkin5 on this equilibrium were recalculated by Lewis and Randall6 and have been reviewed once again in the present work.? The selected values, given in Table 11, agree fairly well with those of Lewis and Randall: The uncertainty in the values is much larger than is indicated by the significant figures recorded, but the relative accuracy is probably greater than the absolute accuracy. TABLE I1 HPDROLYSIS CONSTANT O F CHLORINE K (HOC1) (€If) (C1-)y2tHC!I/(C1Z) Temp., ‘C. K x 104 0.0 1.45 12.8 2.88 25.0 4.47 35.0 5.68
In each experiment there were large changes in the hypochlorous acid (or chlorine) and hydrogen peroxide concentrations, as well as smaller changes in the hydrogen ion and chloride ion concentrations. Thus i t is difficult t o study the concentration dependence for a single substance, but rather all must be considered simultaneously. I n the following analysis the effects have been separated as completely as possible, but it is realized that the best proof of the kq law is the agreement of the data taken as a whole. Hydrogen Peroxide and Hypochlorous Acid Dependence.-It was not possible to vary the initial hydrogen peroxide and hypochlorous acid concentrations greatly, since they had to be small t o give a measurable rate and much further decrease greatly reduced the accuracy of titration. Also, they had t o be of t h e same order of (5) A. A. Jakowkin, 2. physik. Chcm., 29, 613 (1899). (6) G. N. Lewis and M. Randall, “Thermodynamics.” McGrawHill Book Co., New York, N. Y., 1923,p. 508. (7) R. E. Connick, Thesis, University of California, 1942.
INTERACTION OF HYDROGEN PEROXIDE AND HYPOCHLOROUS ACID
June, 1947
1511
magnitude since the sum of the two was measured and the concentration of the less abundant species determined by difference. The hydrogen peroxide dependence is best tested by an experiment in which its initial concentration is less than that of the hypochlorous acid, so that it falls to zero a t completion of the reaction. The data for such a run are given in Table 111. For approximately an eight-fold variation of peroxide concentration the value of kq is essentially constant. Similarly all other experiments of Table I showed a first order hydrogen peroxide dependence. TABLE I11 DEPENDENCE OF RATE ON HYDROGEN PEROXIDE
CON-
CENTRATION
ExDt. 75 V I I ., temD. - 25'. _LC. = 0.0958 Time, mu.
0.483
0.842 1.277 1.87 2.88 3.81
(HrO,) (HOCI) (HCI) X 10' X 10' X 10' moles/liter moles/liter moles/liter
13.72 11.06 8.40 5.75 3.09 1.76
25.70 23.02 20.35 17.71 15.06 13.75
- d(H,Or) df
I
I
-6 kr
X 10'
X 10-7
33.92 36.56 39.21 41.81 44.54 45.89
7.97 6.74 5.36 3.59 1.89 1.02
1.97 1.98 2.04 2.02 2.05 2.00
-5
I
f
-4
-3
I
-2
I
I
r
-1
0
1
Log ( H + ) ( C l - ) . Fig. 1.-kZo as a function of hydrogen and chloride ion concentrations: 0 data from this paper; Q data from Makower and Bray.1
the points fall off to a horizontal line which is Makower and Bray's' limiting law for high hydrochloric acid. Ionic Strength Dependence.-The variation TABLE IV of kr with ionic strength a t low hydrochloric acid DEPENDENCE OF ~ T ON E HYPOCHLOROUS ACID CON- concentration is shown in Table V. The experimental conditions were similar for all runs except CENTRATION for the variation of ionic strength which was alExpt. 96 IV, temp. 2 5 O , = 0.289 tered by the addition of sodium perchlorate (see ( H O ) (HOCI) (C1-1 ( H + ) - d(HnO,) x 10' x 10' x 103 x 10' Table I for typical conditions for run 75 VII). Time, moles/ moles/ moles/ moles/ df kd mtn.
liter
liter
liter
liter
X 10'
X 10-7
0.59 1.04 1.57 1.86
5.62 4.46 3.72 3.46
2.66 1.69 1.06 0.84
7.40 7.54 7.62 7.65
19.1 19.2 19.3 19.3
3.59 1.78 1.02 0.76
1.70 1.63 1.76 1.77
The data of Table IV illustrate the first order dependence of the rate on the hypochlorous acid concentration over a three-fold variation. If the dependence on H + and C1- is assumed to be known, a nearly three-hundred-fold variation of HOC1 concentration may be obtained by comparing experiments 75 VI1 and 107 VI of Table I. The relatively small discrepancy in the k d o values is probably due to an error in the chlorine hydrolysis constant, which as pointed out earlier is subject to considerable uncertainty. Hydrogen Ion and Chloride Ion Dependence. -The product of the hydrogen ion and chloride ion concentrations was varied over a very wide range in the experiments of Table I. The data have been plotted in Fig. 1 using a rate constant, kZ0, defined by the equation
TABLE V
EFFECTOF IONICSTRENGTH ON k,
AT
25"
Expt.
P
kr X 10-7
Extrapolated 89 VI1 86 VI1 75 VI1 80 VI1 84 VI1
0 0.0036
3.13 2.76 2.41 2.00 1.98 2.16
,0242 .0958 .3725 1.015
One might expect the ionic strength effect to be related to the change in the activity coefficient of hydrochloric acid. In Fig. 2, it is seen that this expectation is well fulfilled. The solid line is a plot of the mean activity coefficient of hydrochloric acid against the square root of the ionic strength, which in this case is the square root of the hydrochloric acid concentration. The solid circles are the data tabulated in Table VI the quantity plotted being '/*(log k d - 7.406). One point is placed arbitrarily on the hydrochloric acid curve, the position of the remaining points then being fixed. The factor of one-half is neces- d(H*Oz)/dt = k z D ( H 2 0 z ) ( H O C l ) (10) sary since the logarithm'of the mean activity co(kzo is the value of the rate constant at zero ionic efficient of hydrochloric acid is plotted. All values strength obtained by the method described in the including activity coefficients are based on confollowing section). In agreement with the kd law centration units of moles per liter. the plot of log k 2 O against log (H+)(Cl-) yields a From Fig. 2 it is clear that the main &ect of straight line of unit slope at low hydrochloric ionic strength on k d is equal to the change in the acid concentrations. Above (H+)(Cl-) = 0.1 activity coefficients of hydrogen ion and chloride
ROBERT E. CONNICK
1512
-0.2
I1
I
I
0
0.2
I
0.4
0.6
I
I
0.8
1.0
4G Fig. 2.-Effe:ct of ionic strength on ka: 0 , '/z (log kd 7.496), d , - l / * (2.238 log ks) (see section on hypochlorite reaction), solid curve, log - y * ~ ~ 1 .
+
ion. The point a t .\/F = 0.610 may be in error, but actually deviations of the points from the curve a t high ionic strength would not be surprising since in one case the ionic strength is due to hydrochloric acid and in the other case mainly to sodium perchlorate. Values read from the hydrochloric acid activity coefficient curve were used to correct values of k* t o zero ionic strength, i. e., t o obtain k$. Temperature Dependence of kr.-Values of k4 a t various temperatures are given in Table VI. The conditions were similar for each experiment (see Table I and I11 for details for run 75 VII). Within experimental error the values of k4 lie on a straight line when plotted against the reciprocal of the absolute temperature. TABLE VI EFFECT OF TEMPERATURE ON kr Expt.
101 VI1 106 VI1 75 VI1 109 VI1
Temp.,
0.00 12.80 25.00 35.00
OC.
* 0.05
* *
=t
.05 .05 .05
k4 x 10-7 1.06 1.48
2.00 2.62
A single experiment a t 0" in the presence of a high concentration of hydrochloric acid, which repressed the hydrolysis of chlorine, gave k4O = 1.41 X 10'. This value corresponds to k4 = 0.90 X lo7 at p = 0.0958, the ionic strength used in the experiments of Table VI. The most likely cause of the 15% discrepancy with experiment 101 VI1 is an error in the hydrolysis constant of chlorine which had t o be used to calculate k4 for the experiment a t high hydrochloric acid concentration. Discussion of the kq Law.-The experimental results amply confirm the mechanism proposed by Bray, which was discussed in the Introduction. The extent of the agreement is shown in Fig. 1, where the solid curve was calculated from equation (7) by fitting the curve to the k4 law at
Vol. 69
low values of (H+)(Cl-) and using Makower and Bray's limiting rate at high values of (H+)(Cl-). The curve quite accurately follows the experimental points, even through the transition region. Since no other reasonable mechanism has been found which fits the experimental data, i t is believed that the above mechanism is correct and that chlorine and hydrogen peroxide react t o form the peroxy compound, HOOCI, which then decomposes t o give oxygen, hydrogen ion and chloride ion. The rate determining step depends on the hydrochloric acid concentration, and is the decomposition of HOOCI a t high values of (H+) (C1-) and the reaction of chlorine with hydrogen peroxide a t low values. The experimental results are consistent with the idea that the concentration of HOOCl is always small relative to the hydrogen peroxide and hypochlorous acid (or chlorine) concentrations. According to the above mechanism, variation of the ionic strength can affect kq through changes in K and k,. The principal effect on K is caused by changes in the activity coefficients of hydrogen ion and chloride ion, while there should be little effect on k,, the rate constant for the reaction between two uncharged molecules. Then the change in kq should correspond t o that of the product of the activity coefficients of hydrogen and chloride ions-a result which has already been verified (see Fig. 2). At high concentrations of hydrochloric acid where the rate law is that given by equation ( 2 ) ) the main effect of changing ionic strength should be to alter the value of the equilibrium constant of reaction ( 5 ) . This assumes a negligible effect of ionic strength on the rate of decomposition of the neutral molecule HOOCI. The equilibrium constant of reaction ( 5 ) will be affected principally by the change in the activity coefficients of hydrogen ion and chloride ion, which is just the ionic strength effect found by Makower and Bray.' The value of the specific rate, ka0, for the bimolecular reaction between hydrogen peroxide and chlorine is found to be 1.1 X lo4 liters moles-' min.-l a t 25O, and 2.1 X lo3 a t 0". These data were taken from the experiments a t (H+)(Cl-) = 0.008 in order t o avoid the use of the hydrolysis constant for which the temperature dependence is uncertain. The activation energy, E , defined by the equation k Ze-E/RT (11) (where E and Z are assumed to be temperature independent) is calculated to be 10,700 calories. There are indirect indications that this activation energy is not actually temperature independent. The data of Table VI show that the k., reaction has a constant activation energy of 4,500 calories from 0 to 35'. This value combined with the AH of the hydrolysis reaction should give the activation energy for the bimolecular reaction between chlorine and hydrogen peroxide, calculated above. According to J a k ~ w k i n there ,~ is a large I
INTERACTION OF HYDROGEN PEROXIDE AND HYPOCIILOROUS ACID
June, 1947
variation of A H hydrolysis with temperature and hence, the activation energy for the bimolecular reaction should show a similar large variation with temperature. This would mean that the ' value of 10,700 calories would be too small a t 0 and too large a t 25' and the variation would be of the order of 2000 calories from the mean. But before any significance can be attached to this conclusion, the anomalous variation of the hydrolysis constant with temperature should be checked. Using 10,700 calories for the activation energy one calculates an Arrhenius constant, 2, for equation (11)) of 0.8 X 10l2liters moles-'min.-' a t 2 5 O , which is roughly a factor of twenty smaller than that to be expected from collision theory. On the basis of absolute reaction rate equations, the entropy of activation is -14.6 entropy units when concentrations are expressed in moles per liter. The Hypochlorite Reaction A t very low hydrogen ion concentrations the reaction of hydrogen peroxide with hypochlorous acid becomes much faster than would be predicted from the k4 law, thus indicating the existence of a new reaction path. To study this reaction, i t was necessary to drown out other reaction paths by using very low concentrations of hydrogen ion and chloride ion, The experimental results are given in Table VII. Listed are the initial concentrations of the reactants, the ionic strength p, and two rate constants defined by the equations
-
- d(HzO*)/dt = kz(HzOz)(HOCl) d(HaOz)/dt = kB(HzOz)(HOCI)/(H+)
(12) (13)
1513
18 VII, where the contribution from this extraneous reaction is small, show the reaction to be independent of the chloride ion concentration. Comparison of 15 VI1 and 18 VI1 gives some evidence of first order dependence on hypochlorous acid concentration. I n the first four experiments of Table VI1 the value of kg increased throughout each experiment, since the formation of hydrogen ion caused an increase in the contribution from the extraneous reaction. In the last three experiments hypochlorous acid was used, rather than chlorine, in order to minimize the initial hydrogen ion concentration. Values of kg in the early part of the reaction, where the hydrogen ion concentration was very low, gave quite constant results as shown by the data of Table VIII. The high value a t the start is probably attributable to the uncertainty in measuring the large initial value of - d(HzOz)/ dt. These data illustrate very clearly the inverse first power dependence on hydrogen ion concentration. There is also good evidence for a first power hydrogen peroxide dependence. TABLE VI11 Experiment 60 VII; Temp. 25O; p = 0.0001; initial concentrations: (H202) = 2.31 X M, (HOCI) = 2.67 x 10-4 M,(a-) = -0, (H+) = -0
0.63 2.18 5.45 11.03 19.67
2.123 1.946 1.769 1.591 1.414
2.485 2.308 2.131 1.953 1.776
0.189 .361 .538 .716 .893
0.187 .074 .0404
.0253 .0169
0.68 .59 .58 .58 .60
TABLE VI1 Discussion of the Mechanism of the Hypoexperimental data indiREACTIONchlorite Reaction.-The EXPERIMEXTAL DATAFOR THE HYPOCHLORITE Expt.
13VII 15VII 9VII 18VII 5VIII 65 VI1 119 VI1
(HzOz) (HOCl) ((213 X 104 X lo4 X I O 4
2.52 2.54 2.54 2.54 2.05 2.31
1.97 3.70 2.93 1.92 2.64 5.67
2.08 2.67
(H+) X lo4
p
TEXP.25" 1.97 6.21 0,092 3.70 3.70 .092 2.03 2.03 .092 5.09 1.92 .092 -0 -0 ,092 -0 -0 .0001
TEMP. 0" -0 -0
kz
25 31 43 49
kB
X 101
1.50 1.23 0.94 1.06 0.88 -58
cate t h a t the rate law for the hypochlorite reaction is dt
This is most simply interpreted by the following mechanism KB
0.051
The values for kz and kg were calculated for a time shortly after the start of the reaction. Therefore, one cannot use the initial hydrogen ion concentration of column 5 to calculate one from the other, since the hydrogen ion concentration changes considerably during the reaction. The values of k g in Table VI1 decrease with decreasing acidity but appear to be approaching a value of cu. 0.90 X lo-*. (Experiment 65 VI1 is a t a different ionic strength and should not be compared with the others.) This trend in the value of kg is due to another reaction path which will be discussed later. Experiments 9 VI1 and
+ C10+ HzOz +HzO + + C1HOCl = H +
C10.0001
kB (H~OQ) (HOCl) (H +)
- - d(HzO2) =
kd
0 2
(15)
The rapid equilibrium (14) followed by the rate determining step (15) leads to the rate law given above, with kg = kdKB
(16)
From a comparison of this rate law with the k4 law, it is clear that this reaction path is only important relative to the k d path a t very low hydrogen ion and/or chloride ion concentrations. According to the Bronsted theory of neutral salt effects, k d should vary little with ionic strength; however, K B will vary greatly because of the change in the activity coefficients of H and C10-. Substituting H+ and C1- for H + and C10-, for which there are no data available, one +
1514
Vol. 69
KOBEIITE. CONNICK
predicts that k B should increase with increasing ionic strength in accordance with the change of the activity coefficient of hydrochloric acid The data fw experiments 65 VI1 and 5 VIII, of ionic strength 0.0001 and 0.092, respectively, have been plotted in Fig. 2 and are indicated by the symbol d . The quantity plotted is - l/2 (2.238 f log kB) where 2.238 is used to fix one point on the hydrochloric acid curve while the factor '/* is necessary because y r of hydrochloric acid is plotted. It is seen that the points agree well with the predictions of the Bronsted theory. The values chosen a t zero ionic strength for KE are 5.7 x min.-l a t 25' and 5.0 X min.-l a t 0'. Using the ionization constant of hypochlorous acid, i t is possible to calculate the specific rate of reaction, k d r of C10- and hydrogen peroxide. The ionization constant was taken to be 3.44 X 10" a t 25' and 1.82 X a t 0'. These values were calculated from Ingham and Morrison's8 value a t IS', using the heat of ionization of hypochlorous acidg and Pitzer's'O equation for the change of the ionization constant of acids with temperature. The value of kdo is then equal to 1.7 X lo5at 25' and 2.8 X l o 4a t 0'. Calculation by equation (11) yields an activation energy of 11,600 calories and an Arrhenius constant of 5.4 x 1013. In terms of statistical rate theory the entropy of activation is -6.1 entropy units, when concentrations are expressed in moles per liter. The observed rate law can also be interpreted by the following mechanism
ion and chloride ion concentrations. The rate is not completely reproducible and somewhat more than one mole of hypochlorous acid seems to disappear per mole of hydrogen peroxide reacted. A careful attempt to exclude all impurities failed to eliminate this reaction. It was not possible to formulate a rate law. Some experiments were performed in an acetic acid-sodium acetate buffer mixture. Below 10-5 M hydrogen ion concentration the reaction was mainly due to the hypochlorite path under the conditions used. Above this concentration the rate was greater by a factor of 20 to 30 than under the same conditions in the absence of the buffer. Similar results have been reported by Makower.2 The specific nature of this catalytic action of the acetic acid-sodium acetate buffer is not known. Acknowledgment.-The author wishes to express his great indebtedness t o the late Professor William C. Bray for the constant guidance and inspiration during the course of this work. Summary The kinetics of the reaction between hydrogen peroxide and chlorine to give oxygen and chloride ion was investigated in dilute hydrochloric acid solutions. The present results, when combined with those of Makower and Bray for higher hydrochloric acid concentrations, lead to the following mechanism for the reaction Hz09
+ Clz
ka
H+
+ C1- + HOOCl
kb
HOOCl
k, + 0 2 + H + + C1-
In the region above 1 M hydrochloric acid the decomposition of HOOCl is the rate determining step while a t lower acidity the slow step is the reaction of hydrogen peroxide with chlorine. The second order rate constant, ka, was found to have a value of 1.1 X lo4liters moles-' m h - 1 at 25' and zero ionic strength. The activation energy for the reaction is 10,700 calories. I n the region of hydrochloric acid concentrations of and less, where the chlorine is essentially all hydrolyzed to hypochlorous acid, the mechanism of reaction becomes
Making use of the ionization constant of hydrogen peroxide,ll the activation energy and Arrhenius constant for reaction (18) are calculated to be 7,700 calories and 1.1 X 1015liters moles-' min.-l, respectively. The value of the Arrhenius constant seems too great and, therefore, the hypochlorite mechanism is preferred. Rate of Reaction in Region (H+)(CI--) = to 10-"--In the region of (H+)(Cl-) from t o lod8the reaction is faster than can be accounted for by the k4 and hypochlorite paths. HOCl = H + + C10This additional reaction was difficult to study bekd cause of the large correction for the contributions C10- + H202 +H20 0 2 C1from the k, and hypochlorite paths, which seldom was less than 50%. Several characteristics of the The rate determining step is the reaction of hynew reaction path may be mentioned. The rate drogen peroxide with hypochlorite ion. The secis nearly or completely independent of the hypo- ond order rate constant was found to be 1.7 X lob chlorous acid concentration but first order with liters moles-' min.-' a t 25' and zero ionic strength respect to the peroxide concentration. There is a and the reaction between (210- and hydrogen perless than first power dependence on both hydrogen oxide has an activation energy of 11,600 calories. In the region of (H+)(Cl-) equal to to (8) J. W. Ingham and J. Morrison, J . Chem. Soc , 1200 (1933). the reaction proceeds partly by the above two (9) F. R. Bichowsky and F. D. Rossini, "Thermochemistry of the Chemical Substances," Reinhold Publishing Co., N e w York, N. Y., mechanisms and partly by a third mechanism. 1936. Several characteristics of this last reaction path (10) K.S. Pitzcr, TEIS JOURNAL, 69, 2365 (1937),equation 10,p. are mentioned. 2370. BERKELEY, CALXRORNIA RECEIVED AUGUST21, 1948 (11) R. A. Joyncr, 2. ~)rorg.Chrm.. 77, 108 (1912).
+ +
